Chapter 11.3: Molecular Orbital (MO) Theory and Electron Delocalization

Flashcards from Chapter 11.3 of Chemistry: The Molecular Nature of Matter and Change.

Molecular orbital theory (MO theory)

A model of bonding to explain each molecule’s properties through the use of delocalized electrons in molecular orbitals of given energies and shapes that extend over the whole molecule

• Creates a number of bonding and antibonding orbitals equal to the number of atomic orbitals

Bonding molecular orbital (bonding MO)

This is formed through the overlap of two electron clouds, increasing density between nuclei

• Acts through the amplification of two atomic orbital wave functions, adding two functions

• Lower in energy due to density spread between nuclei across larger volumes, lowering repulsions for a more stable state

• Always coexists with an antibonding MO

Antibonding molecular orbital (antibonding MO)

This is seen in nodes, or regions of zero electron density between the nuclei

• Acts through the cancellation of two atomic orbital wave functions, subtracting one function

• Higher in energy due to greater nuclear repulsions and a lack of separating electrons via a node, lowering stability

• Always coexists with a bonding MO

Sigma molecular orbital (σ MO)

A molecular orbital that is cylindrically symmetrical about the nuclei, formed through similar energy and orientation

• Denoted by a σ with a subscript listing the derived atomic orbital, such as σ_1s

• Antibonding molecular orbitals are denoted with a star, such a σ^*_1s

Molecular orbital diagram (MO diagram)

A diagram showing the atomic orbitals that compose each MO, following normal diagram rules

• Fills lower-energy bonding before higher-energy antibonding forms

• Each level is written in parentheses; for H₂, this is (σ_1s)²

Molecular orbital bond order (MO bond order)

Calculated by taking:

• ½ (number of bonding electrons - number of antibonding electrons)

  • Greater than 0 indicates more stability and formation, increasing with strength

  • Equal to 0 indicates no difference in stability, resulting in no formation

• Correlates with greater bond energy and shorter bond length

Homonuclear diatomic molecule

A molecule composed of two identical atoms; these only use outer atomic orbitals to determine molecular orbitals

Pi molecular orbital (π MO)

A molecular orbital formed through the side combination of two p orbitals with orientations perpendicular to the internuclear axis

• Noted in pairs like π_2p and π^*_2p

• Expected to have a σ_2p < π_2p < π^*_2p < σ^*_2p energy order for the smaller O₂, F₂ and Ne₂

• Expected to have a π_2p < σ_2p < π^*_2p < σ^*_2p energy order for the larger B₂, C₂, and N₂ due to sp mixing raising the energy of the σ_2p and σ^*_2p orbitals

sp Mixing

This occurs for the larger homonuclear diatomic molecules B₂, C₂, and N₂, with energies being lowered on the σ_2s and σ^*_2s and raised on the σ_2p and σ^*_2p orbitals due to less repulsion

• Deviates from the expected non-mixing of two orbitals for normally large energy differences, as seen in O₂, F₂, and Ne₂

• Results in molecular bond energy being listed as π_2p < σ_2p < π^*_2p < σ^*_2p

Magnetism

This can be observed in a molecular orbital diagram, determining if any unpaired electrons are present in a molecular orbital while avoiding the shortcomings of Lewis structures

Heteronuclear diatomic molecules

These have asymmetric MO diagrams due to unequal atomic energies and nuclear charges, with greater charges resulting in more stable forms

Nonbonding molecular orbital (nonbonding MO)

A molecular orbital that is not involved in bonding, with the same energy as the isolated orbital

• Seen with the two filled pairs in F not interacting with H in HF